Honors Unit 3: The Electron and Periodic Table Packet PDF

Name KEY Honors Unit 3 The Electron and Periodic Table Packet Draw the evolution of the atomic model and provide a description next to each drawing: Drawing Description Dalton (1803) Explain why Dalton...

Name KEY Honors Unit 3 The Electron and Periodic Table Packet Draw the evolution of the atomic model and provide a description next to each drawing: Drawing Description Dalton (1803) Explain why Dalton’s model of the atom changed? Discovery of the electron Thomson (1897) Explain why Thomson’s model of the atom changed Discovery of the atomic nucleus Rutherford (1911) Explain why Rutherford’s model of the atom needed modification? Why doesn’t the atom collapse? 1 Bohr (1913) Explain why the Bohr model of the atom had to be modified. Discovery of the neutrons, electrons do not orbit nucleus Quantum Mechanical (1926) Heisenberg – Heisenberg Uncertainty Principle It is impossible to know both the velocity and position of a particle at the same time. Once you know one, you have altered the other. (This applies to small particles, such as electrons) 2 The quantum mechanical model grew out of the study of light. In the beginning, light was thought to consist of particles, but by 1900, scientists learned that light consisted of waves. Properties of waves: Amplitude Height of wave from zero to crest, describes brightness Frequency # of waves to pass given point per unit of time (usually second) Hertz (Hz), s-1 Wavelength Distance between 2 crests Meters (m), nanometers (nm) Wave A Wave B Which wave has a shorter wavelength? Wave A Which wave has a higher frequency? Wave A From your answers above, what is the relationship between frequency and wavelength?Frequency and wavelength are inversely proportional – as one increases, the other decreases 3 Why do certain elements give off certain colors when heated? For example, strontium gives off a red color and copper gives off a blue color. Color is given off when electrons move from an excited state (higher energy) to a ground state (lower energy). All elements have a different combination of neutrons, protons & electrons so each has its own emission spectrum. Scientists realized that distinct lines of colored light produced were produced when electrons went from excited state (higher energy) to ground state (lower energy). The frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element. 4  Atomic Spectra Activity Introduction: Each element has its own unique line spectrum and is thus referred to as the “fingerprint” for a particular element. The spectra for each element are unique because each element contains differing numbers of electrons and thus different energy levels. When energy is absorbed (plugging in the light), the electrons jump up to a higher energy state known as the excited state. This is not a stable state so they “fall” back down to the ground state. This process releases energy in the form of light. Think about the following questions as you look at the spectral lines corresponding to each gas: Are the lines closely packed, or spread out over many different colors? Are there many lines you can see, or only a few? How do the colors of the lines from each tube relate to the color you see from each tube when you don't look through the gratings? Procedure Go through all four stations. Directions posted at stations. Spectra Observations: Element: Tube color - Element: Tube color - 5 Indirect Sunlight: Classroom Lights: Analysis: 1) How are electrons “excited” in this part of the lab? What happens when the electrons relax? 2) What do the different colors in a line spectrum represent? Why are the spectra for each element unique? 3) Which element produced the largest number of lines? Which element produced the smallest number of lines? Explain why elements produce different numbers of spectral lines. 6  Atomic Orbitals An atomic orbital is a region of space in which there is a high probability of finding an electron. These atomic orbitals, or sublevels, can be of different energies and shape. The s orbitals are spherical, have 1 orientation, and can hold max of 2 e- – The p orbitals are dumbbell shaped, have 3 orientations, and can hold max of 6 e- – 7 The d orbitals have 5 different orientations and can hold max of 10 e- – The f orbitals have 7 different orientations and can hold max of 14 e- - 8  Electron Configuration Electron configuration is the way in which electrons are arranged in various orbitals around the nuclei of atoms. Each electron can be represented by three components: Coefficient Represents energy levels Distance from the nucleus Values – 1, 2, 3, 4, 5, 6, 7 Letter Represents the shape of the orbital, sublevel Values – s, p, d, f Superscript Represents the number of electrons in that orbital Values – s = 1 or 2 p = 1, 2, 3, 4, 5, 6 d = 1, 2, 3, 4, 5, 6, 7, 8, 9, 10 f = 1, 2, 3, …, 12, 13, 14 9 Examples: Potassium (K) 4s1 Aluminum (Al) 3p1 Coefficient – 4 Coefficient – 3 Letter – s Letter - p Superscript – 1 Superscript – 1 Zinc (Zn) 3d10 Einsteinium (Es) 5f 10 Coefficient – 3 Coefficient - 5 Letter – d Letter - f Superscript – 10 Superscript – 10 You try: Carbon (C) 2p2 Argon (Ar) 3p6 Coefficient – 2 Coefficient – 3 Letter – p Letter – p Superscript – 2 Superscript – 6 Nickel (Ni) 3d8 Boron (B) 2p1 Coefficient – 3 Coefficient – Letter – d Letter – Superscript – 8 Superscript – * There are two exceptions * s2d4 become s1d5 and s2d9 become s1d10 Copper (Cu) Tungsten (W) Coefficient – 3 Coefficient – 5 Letter – d Letter – d Superscript – 10 Superscript – 5 This happens because an element would rather have half-filled sublevels than no particular order. And by taking 1 e- from the s sublevel and giving it to the d sublevel, this more stable arrangement is achieved. 10 When chemists write the electron configuration for an element, they do not just describe one electron, but instead describe all of the electrons within the element. In order to do this, you must follow three rules: Aufbau Principle Always fill lowest energy orbitals first Pauli Exclusion Principle Each electron has its own spin Hund’s Rule Each orbital gets 1 electron before it gets 2 Example: Magnesium (Mg ) 1s22s22p63s2 Sulfur (S) 1s22s22p63s23p4 Valence shell = 3 Valence shell = 3 Valence electrons = 2 Valence electrons = 6 Notice how they both start with 1s2. You always start filling the electrons at the lowest possible energy level (Aufbau Principle) so therefore every configuration will start with helium’s notation, 1s2, and build up from there. When a sublevel is completely filled, you only write one notation* You try: Argon (Ar) Atomic # 18 Valence shell = 3 1s22s22p63s23p6 Valence electrons = 8 Nickel (Ni) Atomic # 28 Valence shell = 4 1s22s22p63s23p64s23d8 Valence electrons = 2 Boron (B) Atomic # 5 Valence shell = 2 1s22s22p1 Valence electrons = 3 * Check your answers by adding the superscripts, they should equal the atomic number * 11 Write the electron configuration for each of the following: 1) Lithium (atomic # = 3) Valence shell = 2 Valence electrons = 1 1s22s1 2) Arsenic (atomic # = 33 ) Valence shell = 4 Valence electrons = 5 1s22s22p63s23p64s23d104p3 3) Calcium (atomic # = 20 ) Valence shell = 4 Valence electrons = 2 1s22s22p63s23p64s2 4) Phosphorus (atomic # = 15 ) Valence shell = 3 Valence electrons = 5 1s22s22p63s23p3 5) Neon (atomic # = 10 ) Valence shell = 2 Valence electrons = 8 1s22s22p6 6) Mercury (atomic # = 80 ) Valence shell = 6 Valence electrons = 2 1s22s22p63s23p64s23d104p65s24d105p66s24f145d10 12  Noble Gas Notation Instead of writing out the whole electron configuration, we can write in noble gas notation also called short hand notation. Step 1: Find the noble gas that comes before the element you are writing the configuration for Step 2: Write the noble gas in brackets Step 3: Write the electron configuration from the noble gas to the element Element Noble Gas Noble Gas Configuration Aluminum (Al) Neon (Ne) [Ne] 3s 3p1 2 Practice Problems: Element Noble Gas Noble Gas Configuration Check 18 + 2 + 5 = 25 Manganese (Mn) Argon (Ar) [Ar]4s23d5 36 + 2 = 38 Strontium (Sr) Krypton (Kr) [Kr] 5s2 2+2+2=6 Carbon (C) Helium (He) [He]2s22p2 36 + 2 + 10 + 4 = 52 Tellurium (Te) Krypton (Kr) [Kr] 5s24d105p4 * Check your answers by adding the superscripts, they should equal the atomic number * 13 Write the shorthand / Noble Gas configuration for each of the following: 1) Aluminum (atomic # = 13) [Ne]3s23p1 2) Arsenic (atomic # = 33) [Ar]4s23d104p3 3) Calcium (atomic # = 20) [Ar]4s2 4) Phosphorus (atomic # = 15) [Ne]3s23p3 5) Zinc (atomic # = 30) [Ar]4s23d10 6) Gallium (atomic # = 31) [Ar]4s23d104p1 7) Mercury (atomic # = 80) [Xe]6s24f145d10 14  Orbital Notation Chemists’ will also draw orbital notations when describing the electrons in an atom. Dashes Arrows Represent orbital shapes Represent electrons s _ (1) d _ _ _ _ _ (5) p _ _ _ (3) f _ _ _ _ _ _ _ (7) Example: Argon (Ar) 1) First draw and label orbitals (dashes) __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 2) Next draw in electrons (arrows) – remember opposite spins 1s 2s 2p 3s 3p You try: Magnesium Atomic # 12 Valence shell = 3 Valence electrons = 2 Nickel (Ni) Atomic # 28 Valence shell = 4 Valence electrons = 2 * Check your answers by adding the arrows, they should equal the atomic number * 15 Draw the orbital notation for each of the following: 1) Lithium (atomic # = 3) Valence shell = 2 Valence electrons = 1 2) Arsenic (atomic # = 33 ) Valence shell = 4 Valence electrons = 5 3) Calcium (atomic # = 20 ) Valence shell = 4 Valence electrons = 2 4) Phosphorus (atomic # = 15 ) Valence shell = 3 Valence electrons = 5 5) Neon (atomic # = 10 ) Valence shell = 2 Valence electrons = 8 6) Cadmium (atomic # = 48 ) Valence shell = 5 Valence electrons = 2 16  Dot Diagrams Valence Electrons - electrons in the valence shell (involved in bonds) Dot diagrams – shows element symbol with valence electrons Magnesium Atomic # 12 Dot diagram Electron configuration - 1s22s22p63s2 Valence shell = 3 Valence electrons = 2 Group # = 2 Sodium Atomic # 11 Valence shell = 3 Valence electrons = 1 1s22s22p63s1 Sodium Atomic # 11 Dot diagram Electron configuration - 1s22s22p63s1 Valence shell = 3 Valence electrons = 1 Group # = 1 Nitrogen Atomic # 7 Dot diagram Electron configuration - 1s22s22p3 Valence shell = 2 Valence electrons = 5 Group # = 5 * Check your answers by checking the group # for the element * 17 Practice Problems For the following elements: a) write electron configuration b) write noble gas configuration c) draw orbital notation d) identify valence shell e) identify valence electrons f) draw dot diagram 1) Selenium a) 1s22s22p63s23p64s23d104p4 b) [Ar] 4s23d104p4 c) d) 4 e) 6 f) 2) Barium a) 1s22s22p63s23p64s23d104p65s24d105p66s1 b) [Xe] 6s1 c) d) 6 e) 1 f) 3) Explain how you followed Aufbau’s Principle when completing your answers for barium. The electrons filled the lowest energy level before being added to a higher energy level: 1s filled then 2s then 2p then 3s etc. 18 Determine what elements are denoted by the following electron configurations: 4) 1s22s22p63s23p4 Sulfur 5) 1s22s22p63s23p64s23d104p65s1 Rubidium 6) [Kr] 5s24d105p3 Anitmony Determine if the following electron configurations are valid. If it is valid, identify the element. If not, make the needed correction and identify the element. 7) 1s22s22p63s23p64s24d104p5 not valid – should be 3d10 not 4d10 8) 1s22s22p63s33d5 not valid – should be 3s2 not 3s3, then should be 3p5 not 3d5 9) [Ra] 7s2 4f106d10 not valid – should be [Rn] not [Ra], should be 4f14 not 4f10 10) How are the dot diagrams similar for two elements in the same family? All dot diagrams will have same number of valence electrons within the same family. 19  Contributions to the Periodic Table Johann Wolfgang Dobereiner (1829) Published classification system Triads – 3 elements with similar properties Cl (35.45) Br (79.90) I (126.90) Stanislao Cannizzaro (1860) Presented a method for calculating atomic mass Dmitri Mendeleev (1869) First Periodic Table Placed elements in order of atomic mass Left blank spaces & predicted properties of yet to be discovered elements Some elements out of order – Example: Tellurium (#52) and Iodine (#53) He thought the masses were wrong. 20 William Ramsay (1900) Identified noble gases on Periodic Table Isolated neon, krypton, and xenon Henry Moseley (1913) Determined atomic number for each elements Elements arranged by atomic number Worked under Ernest Rutherford Glenn Seaborg (1940) Discovered transuranic elements 94 – 102 Pulled out inner transition from main body of the Periodic Table Actinide placed below lanthanides Periodic Law – When elements are arranged in order of increasing atomic number, there is a periodic repetition of physical and chemical properties. 21 Click on the link, edPuzzle History of the Periodic Table, to answer the following questions: 1) Dmitri Mendeleev organized elements according to 2) This lead to elements being arranged in order of increasing. 3) What did Mendeleev leave in his periodic table? What did he think it would lead to? 4) Why was Mendeleev’s Periodic Table accepted? 5) How was Mendeleev’s problem of some elements being in the “wrong” order of atomic mass fixed? 6) Periods are on the Periodic Table and have the same number of 7) Groups are on the Periodic Table and have similar because they have the same number of electrons. 22  Metals, Metalloids, and Nonmetals Click the link in schoology to watch the video on metals, metalloids, and nonmetals. Characteristics: Metals Metalloids Nonmetals - Good conductors of heat/electricity Properties of both Poor conductors (C is exception) - Luster (shiny) - Solids, except mercury Brittle – shatter if hit - Ductile (pulled into wires) Most gases, some solids and liquids - Malleable (bendable) - 80% of elements gain electrons to form anions - lose electrons to form cations 23 Metals, Metalloids, Nonmetals. Complete the following procedures to determine if your samples are metals, metalloids, or nonmetals. Procedure 1 – Conductivity. Plug in the apparatus. If the light bulb lights, it is a conductor. If the light bulb does not light, it is not a conductor of electricity. Sample Conductor A B C Procedure 2 – Malleability. Touch the samples to determine if you can bend the material. If you can bend it, it is malleable. If you cannot bend it, it is not malleable. Do NOT touch the samples!! Sample Conductor A B C 24 Procedure 3 – Appearance. Look at your samples in the baggies. Record if each has luster (shiny) or is dull. Do NOT touch the samples!!! Sample Conductor A B C Looking over your results, determine if the samples are metals, metalloids, or nonmetals. Provide a reason with your answer for each. Sample A is a because Sample B is a because Sample C is a because 25  Periods and Families Rows on the periodic table are called Periods Columns / Groups on the periodic table are called Families Alkali Metals Alkaline earth metals Extremely reactive metals Reactive metals Group 1 A Group 2A 1 valence electron 2 valence electrons Transition metals Inner Transition Metallic properties Found below main body of PT Charge varies Halogens Noble Gases 7 valence electrons 8 valence electrons Extremely reactive nonmetals Inert – nonreactive Group 7A Group 8 26 We are going to learn about periodic trends. In order to truly understand periodic trends, we must understand a couple of terms/concepts: Nuclear Charge The nuclear charge is the total charge of all the protons in the nucleus. It has the same value as the atomic number. Practice – Provide the nuclear charge for the following elements: 1) Na 11 2) Mg 12 3) Al 13 4) Si 14 5) P 15 6) S 16 7) Cl 17 8) Ar 18 Shielding Electrons in an atom can shield each other from the pull of the nucleus. This effect, called the shielding effect, describes the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. The more electron shells there are, the greater the shielding effect experienced by the outermost electrons. Practice – Circle the atom in each pair that would have the greatest shielding effect: 1) Na Rb 2) Mg Sr 3) Cl Br 4) Si Ge Effective Nuclear Charge, Zeff The effective nuclear charge (often symbolized as Zeff or Z*) is the net positive charge experienced by an electron in a multi-electron atom. The term "effective" is used because the shielding effect of negatively charged electrons prevents higher orbital electrons from experiencing the full nuclear charge. The larger effective nuclear charge results in electrons feeling a greater attraction. This results in smaller size, greater electron affinity, and a larger ionization energy. Practice – Calculate the effective nuclear charge for each atom. (It is the same as the valence electrons) 1) Na 1 2) Mg 2 3) Al 3 4) Si 4 5) P 5 4) S 6 27  Trends Atomic Radius: determined as the distance between the nuclei of two identical atoms bonded together Ionization Energy: amount of energy needed to remove an electron Electronegativity: the attraction on electrons within a chemical bond Electron Affinity: amount of energy released when an electron is added Looking at the definitions above, circle the correct italicized words: 1) If the size of an atom increases, it is harder / easier to remove an electron. 2) If the size of an atom increases, the atom will have a stronger / weaker hold on its own electrons. 3) If the size of an atom increases, the atom will have a stronger / weaker attraction for another atoms electrons. To summarize, the LARGER the atomic radius, the SMALLER the ionization energy, the SMALLER the electronegativity, and the SMALLER the electron affinity. We have group trends and periodic trends but as long as you remember the relationships above, understanding the group and period trends can be easy. 28  Group Trends Complete the following table for four elements in the Alkali family. Name Electron Configuration Valence Shell Lithium 1s22s1 2 Sodium 1s22s22p63s1 3 Potassium 1s22s22p63s23p64s1 4 Rubidium 1s22s22p63s23p64s23d104p65s1 5 Atomic radius INCREASES going down a family on the periodic table because additional ENERGY LEVELS increase the distance between the nucleus and the outermost shell (valence shell). The larger the atom, the EASIER it is to remove an electron. As atomic radius INCREASES , ionization energy DECREASES. With more energy levels, the outermost electrons (the valence electrons) are further from the nucleus and are not so strongly attracted to the nucleus. Thus the ionization energy of the elements decreases as you go down the periodic table because it requires less energy to remove the electrons. The larger the atom, the SMALLER the attraction for the shared electrons. As atomic radius INCREASES , electronegativity DECREASES. The larger the atom, the SMALLER the attraction for the incoming electrons. As atomic radius INCREASES, electron affinity DECREASES. 29 Practice Problems: Determine which atom has a larger atomic radius in each pair and write your reasoning: 1) Aluminum vs. Boron Aluminum, it has an additional energy level. 2) Sodium vs. Rubidium Rubidium, it has additional energy levels. Determine which atom has a larger ionization energy in each pair and write your reasoning: 3) Iodine vs. Chlorine Chlorine, it has a smaller atomic radius so harder to remove electrons. 4) Aluminum vs. Thallium Aluminum, it has a smaller atomic radius so harder to remove electrons. Determine which atom has a larger electronegativity in each pair and write your reasoning: 5) Rubidium vs. Lithium Lithium, smaller atomic radius so stronger attraction to electrons in a bond. 6) Tin vs. Carbon Carbon, smaller atomic radius so stronger attraction to electrons in a bond. Determine which atom has a larger electron affinity in each pair and write your reasoning: 7) Gallium vs. Boron Boron, the smaller the atom, the larger the attraction for incoming electrons. 8) Strontium vs. Calcium Calcium, the smaller the atom, the larger the attraction for incoming electrons. 30  Period Trends Lithium Beryllium Boron Carbon Nitrogen Noble gas [He]2s1 [He]2s2 [He]2s22p1 [He]2s22p2 [He]2s22p3 configuration # Protons 3 4 5 6 7 Effective Nuclear Charge 1 2 3 4 5 Zeff (valence e-) As you move across a period, what happens to the … # of protons: increases by 1 Effective nuclear charge: increases by 1 Atomic radius DECREASES going across a period on the periodic table because EFFECTIVE NUCLEAR CHARGE increase making the electron cloud smaller. The smaller the atom, the HARDER it is to remove an electron. As atomic radius decreases, ionization energy increases. As the nuclear charge increases, the attraction between the nucleus and the electrons increases and it requires more energy to remove the outermost electron. The smaller the atom, the greater the attraction for the shared electrons. As atomic radius decreases, electronegativity increases. 31 The smaller the atom, the greater the attraction for the incoming electrons. As atomic radius decreases, electron affinity increases. Practice Problems: Determine which atom has a larger atomic radius in each pair and write your reasoning: 1) Aluminum vs. Chlorine Aluminum, Chlorine has a higher effective nuclear charge making the electron cloud smaller. 2) Tin vs. Rubidium Rubidium, Tin has a higher effective nuclear charge making the electron cloud smaller. Determine which atom has a larger ionization energy in each pair and write your reasoning: 3) Gallium vs. Calcium Gallium, it is a smaller atom so harder to remove electrons. 4) Cesium vs. Barium Barium, it is a smaller atom so harder to remove electrons. Determine which atom has a larger electronegativity in each pair and write your reasoning: 5) Potassium vs. Bromine Bromine, it is a smaller atom so stronger attraction for electrons in a bond. 6) Fluorine vs. Beryllium Fluorine, it is a smaller atom so stronger attraction for electrons in a bond. Determine which atom has a larger electron affinity in each pair and write your reasoning: 7) Strontium vs. Tellurium Tellurium, it is a smaller atom so stronger attraction for incoming electrons. 8) Bromine vs. Germanium Bromine, it is a smaller atom so stronger attraction for incoming electrons. 32  Ionic Radius Atom vs. Cation Trend A cation is a positive ion and occurs when an atom loses electrons. Write the electron configuration for the calcium atom: 1s22s22p63s23p64s2 Write the electron configuration for the calcium cation (Ca+2): 1s22s22p63s23p6 A cation is smaller than its atom because it has lost electrons, losing an entire energy level. Therefore, effective nuclear charge affecting the remaining electrons increases , pulling electrons in more closely. This is due to a less shielding. Atom vs. Anion Trend An anion is a negative ion and occurs when an atom gains electrons. Write the electron configuration for the sulfur atom: 1s22s22p63s23p4 Write the electron configuration for the sulfide anion (S-2): 1s22s22p63s23p6 An anion is larger than its atom because it has gained electrons, creating more electron-electron repulsions, making the electrons spread out more. 33 Practice Problems: (1) Circle the atom in each pair that has a larger ionic radius (2) Provide an explanation: gains electrons creating more electron-electron repulsions or lost electrons, losing an energy level increasing effective nuclear charge. 1) Bromine vs. Bromide anion gains electrons creating more electron-electron repulsions 2) Sodium vs. Sodium cation lost electrons, losing an energy level increasing effective nuclear charge. 3) Potassium vs. Potassium ion lost electrons, losing an energy level increasing effective nuclear charge. 4) Iodine vs. Iodide ion gains electrons creating more electron-electron repulsions 5) Calcium vs. Calcium ion lost electrons, losing an energy level increasing effective nuclear charge. 34 Test Review For the following elements, (a) write the electron configuration (b) write the noble gas configuration (c) draw the orbital notation (d) determine the valence shell (e) determine the valence electrons (f) draw a dot diagram 1) Sodium (a) 1s22s22p63s1 (b) [Ne] 3s1 (c) (d) 3 (e) 1 (f) 2) Bromine (a) 1s22s22p63s23p64s23d104p5 (b) [Ar] 4s23d104p5 (c) (d) 4 (e) 7 35 (f) 3) Oxygen (a) 1s22s22p4 (b) [He] 2s22p4 (c) (d) 2 (e) 6 (f) Describe the three rules that govern electron configuration. 4) Aufbau Principle: electrons fill lowest energy level first 5) Hund’s Rule: each orbital gets one electron before it gets two 6) Pauli’s Exclusion Principle: no two electrons will have the same spin 36 Label the following families on the periodic table and give two characteristics of each: 7) Alkali metals 9) Halogens reactive, soft metallic solids, shiny, lustrous reactive nonmetals, change state as move melting and boiling points increase with increasing atomic number. 8) Alkaline earth metals 10) Noble gases metallic solids, harder and higher inert, gases melting points than alkali metals, shiny, Label metals, metalloids, and nonmetals on the following periodic table. 11) Metals 12) Metalloids: 13) Non-metals Malleable, ductile, good conductors Properties of both Solids, liquids and 37 of heat and electricity, solids gases, Brittle except for mercury, luster, high melting and boiling point Label the periods on the following periodic table. Label the trend for atomic radius on the following periodic table. Atomic radius – size of an atom 14) Explain why strontium has a larger atomic radius compared to magnesium. Strontium has more energy levels, increasing the distance between nucleus and valence shell. 38 15) Explain why calcium has a larger atomic radius than bromine. Bromine has a larger nuclear attraction pulling the electrons in closer making the atom smaller. Anion – negatively charged ion, gained electrons Cation – positively charged ion, lost electrons 16) Explain why the sulfide anion is larger than the sulfur atom. Sulfide has gained electrons, creating more electron-electron repulsions. 17) Explain why the potassium cation is smaller than the potassium atom. Potassium lost one electron, losing an entire energy level and increasing effective nuclear charge. Label the trend for ionization energy on the following periodic table. Ionization energy – energy needed to remove an electron 18) Explain why Strontium has a smaller ionization energy compared to magnesium Strontium is larger so easier (requires less energy) to remove an electron. 39 19) Explain why bromine has a larger ionization energy compared to calcium. Bromine is smaller so harder (requires more energy) to remove an electron. 20. Write the letters given in the following periodic table to answer the questions. a) Elements in the halogen family? K, P b) Which element in Group IIIA (3A) has the largest atomic radius? S c) Elements with only 2 electrons in valence shell? E, J d) Element with a complete 5d subshell? P e) Element whose last electron is described by 2p3? G f) Element in period 5 with highest ionization energy? L Scientists. Next to each name, write the contribution of each scientist. 21) Mendeleev First Periodic Table, Put elements in order of atomic mass, left blank spaces and predicted properties 22) Moseley Put elements in order of atomic number 23) Seaborg Pulled out inner transition family from main body of Periodic Table 40 24) Cannizzaro Calculated atomic mass 25) Ramsay Added noble gas family 26) You are given samples of 9 elements in sealed glass containers. You have created a table with data about each element. Answer the following questions pertaining to samples 1-9. Element Appearance Additional Information 1 Pale yellow gas Poisonous, reactive element; found in a compound used in toothpaste 2 Silvery solid Very malleable; easily cut with a knife; catches fire spontaneously in water 3 Silvery solid Very malleable; easily cut with a knife; reacts violently with water; found in table salt 4 Colorless gas Inert; used in incandescent light bulbs 5 Silvery solid Fairly hard; found in upset stomach remedies; burns with a bright light in air 6 Colorless gas Inert gas; one of the heaviest gases; used in stroboscopic lamps 7 Silvery solid Rather hard; compounds found in bones and hard water 8 Colorless gas Inert gas; used to fill balloons 9 Greenish gas Poisonous; found in bleach a) Which elements are metals? Write the element number. 2, 3, 5, 7 b) List the elements that belong in the a. Halogen group 1, 9 b. Noble gases 4, 6, 8 c) Do elements 1 and 8 belong in the same group? Why or why not. No because 1 is reactive and 8 is inert. d) Match the element number 3, 6, 7 & 9 from the table with the correct element listed below. Na - 3 Ca - 7 Cl - 9 Xe - 6 41 42 43